[BreakingBad20]

Calculate the standard Gibbs energy and the equilibrium constant of the following reaction at 25°C:      
         Sn(s) + 2AgCl(s)    SnCl2(aq) + 2Ag(s)       

Given the standard reduction potentials of AgCl/Ag as being +0.22V and that of Sn2+/Sn as being -0.14V.

Equilibrium Equation: Kc = ([SnCl2][Ag]^2)/([Sn][AgCl]^2)

Half Equations:
Right: Sn²⁺ + 2e^-  Sn
Left: AgCl  Ag + Cl^-

E° = E°(Right) - E°(Left)
E° = (-0.14) - (+0.22)
E° = -0.36V

I have 2 questions:
Do I use the equation ΔG°cell = −nFE°cel for standard Gibbs energy?

How do I answer this question further?

[BreakingBad20]

Actually I've solved it